Introduction To Redox Types of Reaction Oxidation-Reduction called Redox Ionic compounds are formed through the transfer of electrons. An Oxidation-reduction reaction involves the transfer of electrons. We need a way of keeping track. Oxidation States A way of keeping track of the electrons. Not necessarily true of what is in nature, but it works. need the rules for assigning (memorize). The oxidation state of elements in their standard states is zero. Oxidation state for monoatomic ions are the same as their charge. states state of Oxygen isOxidation assigned an oxidation 2 in its covalent compounds except as a peroxide. In compounds with nonmetals hydrogen is assigned the oxidation state +1. In its compounds fluorine is always –1. The sum of the oxidation states must be zero in compounds or equal the charge of the ion. Oxidation States Assign the oxidation states to each element in the following. CO2 NO3H2SO4 Fe2O3 Fe3O4 Oxidation-Reduction Transfer electrons, so the oxidation states change. Na + 2Cl2 2NaCl CH4 + 2O2 CO2 + 2H2O Oxidation is the loss of electrons. Reduction is the gain of electrons. OIL RIG LEO GER Oxidation-Reduction Oxidation means an increase in oxidation state - lose electrons. Reduction means a decrease in oxidation state - gain electrons. The substance that is oxidized is called the reducing agent. The substance that is reduced is called the oxidizing agent. Redox Reactions Agents Oxidizing agent gets reduced. Gains electrons. More negative oxidation state. Reducing agent gets oxidized. Loses electrons. More positive oxidation state. Identify the Oxidizing agent Reducing agent Substance oxidized Substance reduced in the following reactions Fe (s) + O2(g) Fe2O3(s) Fe2O3(s)+ 3 CO(g) 2 Fe(l) + 3 CO2(g) SO3- + H+ + MnO4- SO4- + H2O + Mn+2 Half-Reactions All redox reactions can be thought of as happening in two halves. One produces electrons - Oxidation half. The other requires electrons - Reduction half. Write the half reactions for the following. Na + Cl2 Na+ + ClSO3- + H+ + MnO4- SO4- + H2O + Mn+2 Balancing Redox Equations In aqueous solutions the key is the number of electrons produced must be the same as those required. For reactions in acidic solution an 8 step procedure. Write separate half reactions For each half reaction balance all reactants except H and O Balance O using H2O Acidic Solution Balance H using H+ Balance charge using e Multiply equations to make electrons equal Add equations and cancel identical species Check that charges and elements are balanced. Practice The following reactions occur in aqueous solution. Balance them Cr(OH)3 + OCl- + OH- -2 CrO4 + Cl + H2O MnO4- + Fe+2 Mn+2 + Fe+3 Cu + NO3- Cu+2 + NO(g) Pb + PbO2 + SO4-2 PbSO4 Mn+2 + NaBiO3 Bi+3 + MnO4- Now for a tough one -4 Fe(CN) + MnO 6 NO3- 4 Mn+2 + Fe+3 + CO2 + Basic Solution Do everything you would with acid, but add one more step. Add enough OH- to both sides to neutralize the H+ CrI3 + Cl2 CrO4- + IO4- + ClFe(OH) + H O Fe(OH) 2 2 2 Redox Titrations Same as any other titration. the permanganate ion is used often because it is its own indicator. MnO4- is purple, Mn+2 is colorless. When reaction solution remains clear, MnO4- is gone. Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect. Example The iron content of iron ore can be determined by titration with standard KMnO4 solution. The iron ore is dissolved in excess HCl, and the iron reduced to Fe+2 ions. This solution is then titrated with KMnO4 solution, producing Fe+3 and Mn+2 ions in acidic solution. If it requires 41.95 mL of 0.205 M KMnO4 to titrate a solution made with 0.6128 g of iron ore, what percent of the ore was iron?